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Electrons and the EM Spectrum. Let’s light stuff on fire! . Models of the Atom.

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Electrons and the em spectrum

Electrons and the EM Spectrum

Let s light stuff on fire

Let’s light stuff on fire! 

Models of the atom

Models of the Atom

So far, the model of the atom consists of protons and neutrons making up a nucleus surrounded by electrons. After performing the gold foil experiment, Rutherford hypothesized a model of the atom that looked much like the one below.

What about the electrons

What about the electrons??

  • Rutherford didn’t know exactly where the electrons were located in the atom, just that they surrounded it, or why chemical bonding occurred.

  • Bohr figured out that electrons orbit in energy levels around the atom.

The bohr model

The Bohr Model

  • Niels Bohr (1852-1962) was a student of Rutherford and believed the model needed improvement.

  • Bohr proposed that an

    electron is found only in

    specific circular paths, or

    orbits, around the nucleus.

    (write this in before “However, this model….”)

Models of the atom1

Models of the Atom

However, this model could not explain the chemical and physical properties of the elements.

Models of the atom2

Models of the Atom

For example it could not explain:

1. why metals give off certain colors when heated in a flame -or-

2 .why objects heated to high temperatures first glow dull red, then yellow, then white

So what is light

So what is light?

Light is all about the electrons

Light is all about the Electrons!

  • Light is a form of electromagnetic radiation that travels like a wave through space as PHOTONS.

  • When electrons get excited, they jump up to higher energy levels and then fall back down

  • Depending on how high they jump, they will give off a different color of LIGHT

Atomic emission spectra

Atomic Emission Spectra

excited state



GAIN energy

LOSE energy

ground state

Energy of electrons

Energy of Electrons

  • When atoms are heated, bright lines appear called line spectra

  • An electron absorbs energy to “jump” to a higher energy level. (excited)

  • When an electron falls to a lower energy level, energy is emitted. (ground level state)

Emission spectrum

Emission Spectrum

Atomic emission spectra1

Atomic Emission Spectra

  • Every element has a UNIQUE emission spectrum. The colors that you see represent the element’s electrons jumping through the energy levels!

  • LONG JUMPS are represented by HIGH energy colors (violet & blue)

  • SHORT JUMPS are represented by LOW energy colors (red).

Em spectrum

EM Spectrum





C. Johannesson

Em spectrum1















EM Spectrum





Properties of light

Properties of Light

  • Movement of excited electrons to lower energy levels, and the subsequent release of energy, is seen as light! Before 1900, scientists thought light behaved solely as a wave. This belief changed when it was later discovered that light also has particle-like characteristics. This is called the wave-particle duality of light.

First let s look at the wave nature of light

First, Let’s look at the WAVE nature of light

Light as a WAVE

Wave properties




Wave Properties

Lesser frequency

greater frequency

Properties of light1

Properties of Light

  • The significant feature of wave motion is its repetitive nature, which can be characterized by the measurable properties of wavelength and frequency.



  • Wavelength () - length of one complete wave

  • Frequency (f) - # of waves that pass a point during a certain time period

    • hertz (Hz) = 1/s

Properties of light2

Properties of Light

  • Electromagnetic radiation is a form of energy that exhibits wavelike behavior (wavelength, frequency, ect) as it travels through space. Together all forms of electromagnetic radiation form the electromagnetic spectrum.

Electrons and the em spectrum

On the electromagnetic spectrum, the lowest energy waves (longest wavelength and lowest frequency) are radio waves. The highest energy waves (shortest wavelength and highest frequency) are gamma rays.

Electrons and the em spectrum

  • FREQUENCY and WAVELENGTH are INVERSELY proportional. (f ↑ ↓)

  • ENERGY and FREQUENCY are DIRECTLY proportional. (E↑ f ↑)

Now let s look at the particle nature of light

Now, let’s look at the PARTICLE nature of light

Light as a PARTICLE

Properties of light3

Properties of Light

  • In the early 1900’s, scientists conducted experiments involving interactions of light and matter that could not be explained by the wave theory of light.

Properties of light4

Properties of Light

  • One experiment involved a phenomenon known as the photoelectric effect.The discovery of the photoelectric effect led to the description of light as having both wave and particle properties.

Electrons and the em spectrum

  • A Quantum of light energy is called a PHOTON.

Em spectrum2

EM Spectrum

  • Frequency & wavelength are inversely proportional

c = f

c:speed of light (3.00  108 m/s)

:wavelength (m, nm, etc.)

f:frequency (Hz)

Example 1

Example 1

If the frequency of a wave is 500 hz, what is the wavelength?

C = λf

C = 3.0 x 108 m/s

f = 500 hz

λ = ?

3.0 x 108 = 500 x λ λ = 600,000 m

One sig fig or 6 x 105

Quantum theory

Quantum Theory

  • The energy of a photon is proportional to its frequency.

E:energy (J, joules)

h:Planck’s constant (6.626  10-34 J·s)

f:frequency (Hz)

E = hf

Example 2

Example 2

What is the energy of a wave if the frequency is 300. hz?

E = hf

f= 300. hz

h = 6.626 x 10-34

E = 300. x 6.626 x 10-34

E = 1.99 x 10-31 3 sig figs

Example 3

Example 3

  • If the energy of a wave is 9.00 x 10-19 J, find frequency and wavelength

    E = hf

    E= 9.00 x 10-19hz

    h = 6.626 x 10-34

    9.00x 10-19 = 6.626 x 10-34 x f

    f = 1.36 x 1015hz 3 sig figs

If the energy of a wave is 9 00 x 10 19 j find frequency and wavelength

If the energy of a wave is 9.00 x 10-19 J, find frequency and wavelength

C = λf

If f = 1.36 x 1015 hz

C = 3.00 x 108 m/s

3.00 x 108 m/s = λ x 1.36 x 1015

λ = 2.21 x 10 -7m 3 sig figs.

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